>>>What is crude? – the simple definition
Crude is a naturally occurring, energy dense hydrocarbon compound that can be extracted (a process mislabeled ‘production’) from subterranean reservoirs and refined into useful products including fuels (including methane, propane, butane, gasoline, naphtha, kerosene, gas-oil, residual oil, and coke), asphalt, and chemical feedstocks. Understanding the chemistry of oil is important to understanding well productivity, production processes and related costs; refining processes and related costs; storage and delivery requirements; product pricing; and environmental impacts. Crude oil, however, is far from homogenous, and understanding the chemistry of crude oil, oilfield geology, and reservoir mechanics provide foundational concepts for further analysis of refining processes and product pricing.
What are hydrocarbons? Pure hydrocarbons are simply molecules comprised of a combination of two elements – hydrogen and carbon. Often other molecules (especially water, salt, metals, carbon dioxide, and hydrogen sulfide) and other elements (especially sulfur and nitrogen) are mixed in with crude. While these ‘impurities’ are ever-present in crude, they are not hydrocarbons in themselves and should be discussed only after providing an overview of hydrocarbon chemistry.
Hydrocarbons are described by their molecular structure – the specific number and combination of hydrogen and carbon atoms. Hydrogen has a valence of 1 and carbon has a valence of 4. As a consequence, hydrogen has a propensity to bond with one other atom, and carbon has a propensity to bond with four. In some instances, however, these bonds can be ‘shared’, in which case we see double or triple bonds between two carbon atoms. This occurs when there are too few hydrogen atoms to satisfy the valence requirements. From a theoretical perspective, these properties infer that an almost infinite number of hydrocarbons can exist. In reality – or should I say in nature or naturally – crude oil is comprised of a mixture of 100,000 to 1,000,000 different types of hydrocarbons along with other elements and other non-hydrocarbon compounds (Raymond and Leffler 2006).
Hydrocarbons are usefully classified into one of two very basic categories: alaphatics and aromatics. Aromatic hydrocarbons contain benzene molecules, but alaphatics do not. The inclusion of benzene is important because it is a known carcinogen, and is therefore regulated by the Environmental Protection Agency. Alaphatic hydrocarbons are further classified by their molecular structure into three categories: paraffins (alkanes), olefins, and naphthenes. As we will see, though, olefins are very unstable, and therefore they tend to bond to other elements also found in reservoirs. For this reason, olefins are rarely found in a crude oil reservoir. Instead they are created by refiners seeking to capture their unique qualities, which are discussed in the section on refining.
Paraffins can be either straight-chained (normal) or branch-chained (isomers), and they can be of any length. Paraffins are considered ‘stable’ hydrocarbons because all the valence requirements are fulfilled with single bonds. Under normal atmospheric temperatures and pressures, paraffins can assume a gaseous, liquid, or solid state depending on the number of carbon atoms per hydrocarbon molecule. These phase characteristics impact production processes and costs, refining processes and costs, and transportation feasibility and costs.
With only one carbon atom and four hydrogen elements, methane (CH4) is the most simple of the paraffinic hydrocarbons, and indeed is the most simple of all hydrocarbons. Ethane (C2H6), and propane (C3H8), are also normal paraffinic hydrocarbons, as is normal butane (nC4H10). Isobutane (iC4H10), on the other hand, is a branch-chained hydrocarbon – the smallest of the branch-chained hydrocarbons, in fact. Functionally, butane and isobutane have the same heat value and are of equivalent thermal efficiency, but their different structures make them behave differently. They boil at different temperatures; they have different specific gravities (thought their molecular gravity is identical); and they can cause different chemical reactions.
Under normal atmospheric pressure and temperature conditions, hydrocarbons with 1 to 4 carbon atoms (methane, ethane, propane, normal butane, and isobutane) exist in a gaseous state, but under higher pressure and/or lower temperatures, these hydrocarbons will undergo a phase change to a fluid state. This is because liquids vaporize when their vapor pressure comes to parity with the surrounding pressure. When vapor pressure is plotted on a pressure-temperature diagram, we see that the relationship is non-linear – with the vapor pressure increasing logarithmically with temperature. The slope of the phase change line (vapor pressure line) is described by the Clausius-Clapeyron relation.
The atmospheric pressure boiling point, or normal boiling point, then, is defined as the temperature at which a liquid (in this case a liquid hydrocarbon) vaporizes under normal atmospheric pressure. The larger the hydrocarbon molecule, the higher the normal boiling point. In descending order the boiling points for the four gaseous hydrocarbons are: 31ºF (butane), -44ºF (propane), -127ºF (ethane), and -258ºF (methane).
Hence we see that under normal atmospheric pressure (1 atm), cooling butane to just below freezing (31ºF) or methane to -258ºF (well below freezing) will cause the gases to undergo a phase change and convert to liquid. Similarly, increasing the surrounding pressure of butane to 2.6 atm, or propane to 14 atm will force the same phase change (provided the temperature does not exceed the new boiling point: 100.4ºF).
When temperatures exceed a compound’s critical temperature, a gas cannot be liquefied by pressure, and below a critical pressure, no amount of cooling can cause a phase change from gas to liquid. The critical temperature of methane is -116.9ºF, hence methane cannot be liquified using pressure alone at normal atmospheric temperatures. The critical temperature of ethane is 89.9ºF, hence under normal atmospheric temperatures, methane and ethane would need to be put under extremely high pressure to liquefy, and methane would have to be kept at very low temperature as well. Both butane and propane, on the other hand, can be easily and economically transformed into liquid form for transportation and storage because these hydrocarbons have a low critical pressure.
Because of their relatively high critical temperatures (relative to normal atmospheric temperatures), methane and ethane are liquefied by a combination of super-cooling and pressurization rather than through compression alone. This is important because super-cooling large amounts of methane and ethane is expensive and energy intensive, though maintaining the sub-boiling point temperature is less so. According to Dr. John Barklay, expert in stranded methane capture and transport, roughly 10% of the total energy content of methane is required to compress and purify methane from stranded fields. While the energy return for energy invested in cryogenic capture and transport of stranded methane has not been researched, it is clearly below 10-to-1.
Paraffins with four to eighteen carbon atoms, assume a fluid state under normal atmospheric temperature and pressure conditions. Alternatively we can say that the normal vapor pressure for these longer paraffins is less than 1 atm, and just as the larger, more complex gaseous hydrocarbons have a higher vapor pressure than their more simple cousins, so it is also true that the larger the liquid paraffins, the higher the pour point (the lowest temperature at which a hydrocarbon assumes a fluid state).
Paraffins with more than 4 carbon atoms can assume branched or straight molecular structures. Those with branched structures have a higher specific gravity. Gasoline is a mixture of numerous hydrocarbons, but the majority of the volume comes from paraffinic alkanes with 5 to 8 carbon atoms each. Normal alkanes produce gasoline with a lower octane number while branched-chain paraffins produce a higher octane number. Alkanes with 9 to 16 carbon atoms are easily refined into diesel and kerosene (jet fuel).
Paraffinic hydrocarbons with more than 18 carbon atoms assume a solid or semi-solid state at normal atmospheric temperatures and pressures. In other words, the pour point, which ranges from 116.6ºF to 147.2ºF for paraffins of these lengths, is higher than normal atmospheric temperature. Hence, under normal atmospheric conditions, these so-called waxy paraffins ‘crystallize’ – which is problematic because the crystallization of waxy paraffins clogs well tubing and reduces wellhead production rates much in the same way that cholesterol crystallizes as plaque on arterial walls restricting blood flow. Wax deposits at best cause production to slow, but at worst, the build up of paraffins in the well tubing causes wells to be shut in so that the deposits can be removed. Alkanes with 17 to 25 carbon atoms are easily refined into fuel oils and lubricating oils.
Olefinic hydrocarbons are unsaturated – meaning that some of the carbon bonds are shared. Unsaturated hydrocarbons are unstable, or volatile, because C=C double bonds and C≣C triple bonds are attracted to and easily bond with other ‘available’ atoms (especially sulfur and nitrogen). The most common of the olefinic hydrocarbons are ethylene and propylene, which are both mono-olefins (olefins which share one single double C=C bond). Diolefins have at least two C=C double bonds, and alkynes have at least one C≣C triple bond. Olefinic hydrocarbons could in theory exist in nature, but for the most part they do not precisely because they are so volatile. Olefins, instead, are produced in refineries to either be used as important feedstocks for the petrochemical industry, or to be used as inputs to the alkylation process (a refining process in which olefins are bonded with isobutane to create a branched-chain isoparaffin called alkylate which is blended with gasoline in order to increase the octane number).
Naphthenic hydrocarbons are similar to paraffins in that they are saturated and, therefore, stable elements. Naphthenic crude tends to be heavier than paraffinic or waxy crude because the ratio of the heavier carbon atoms to lighter hydrogen atoms is greater. This is the case because naphthenic hydrocarbons exist as closed rings rather than straight chains. Like paraffins, naphthenic hydrocarbons can, and often do, have branches which form isomers. If this is the case, the hydrocarbon molecules will find it more difficult to become excited, and the critical temperature will be relatively high.
Aromatic hydrocarbons are thus named because they emit what can be described as a sickly sweetish scent. In part aromatics are predisposed to being fragrant because they are unstable. They have too few hydrogen atoms to satisfy all of the valence requirements. By definition, aromatic hydrocarbons must also contain at least one ring of benzene, which as we shall see in the section on refining, makes them a valuable feedstock because they yield high octane ratings, but also makes them more expensive to produce and handle. While the benzene content of aromatic hydrocarbons increases a gasoline’s octane ratings (as do tuolene and xylene – which are often referred to in the aggregate as the BTX family), aromatic hydrocarbons when refined produce relatively small quantities of gasoline, and relatively high quantities of residual fuel (commonly referred to as ‘resid’).
Asphaltenes are sometimes considered in their own category, but are often considered a subset of aromatic hydrocarbons because they contain benzene. Asphaltenes are often referred to as bitumen, and they are very dense. They have the highest carbon-to-hydrogen ratios of all commonly found hydrocarbons. Asphaltenes appear to be black or tar-like precisely because they are so dense. Asphaltic crude is commonly referred to as ‘heavy crude’ because they typically contain 40 to 70 carbon atoms arranged into complicated branching and cyclic structures. Asphaltenes have been described as “structures that look like balled up chicken wire.” (Raymond and Leffler), and their structure means that under normal atmospheric temperatures and pressures, they exist in a solid to semisolid state. Bitumen can form in one of two ways. If exposed to the atmosphere, lighter hydrocarbons will evaporate or be consumed through bacterial activity leaving only the heaviest, most complex structures. Bitumen can also result from incomplete transformation of organic material to hydrocarbons (see discussion on hydrocarbon formation). Because asphaltenes have such high pour points, they must be heated or chemically treated in order to be produced and transported.